Article pubs.acs.org/est
Control of Heterogeneous Fe(III) (Hydr)oxide Nucleation and Growth by Interfacial Energies and Local Saturations Yandi Hu,†,§ Chelsea Neil,† Byeongdu Lee,‡ and Young-Shin Jun*,† †
Department of Energy, Environmental & Chemical Engineering, Washington University, St. Louis, Missouri 63130, United States X-ray Science Division, Argonne National Laboratory, Argonne, Illinois 60439, United States
‡
S Supporting Information *
ABSTRACT: To predict the fate of aqueous pollutants, a better understanding of heterogeneous Fe(III) (hydr)oxide nucleation and growth on abundant mineral surfaces is needed. In this study, we measured in situ heterogeneous Fe(III) (hydr)oxide nucleation and growth on quartz, muscovite, and corundum (Al2O3) in 10−4 M Fe(III) solution (in 10 mM NaNO3 at pH = 3.7 ± 0.2) using grazing incidence small-angle X-ray scattering (GISAXS). Interestingly, both the fastest heterogeneous nucleation and slowest growth occurred on corundum. To elucidate the mechanisms, zeta potential and water contact angle measurements were conducted. Electrostatic forces between the charged Fe(III) (hydr)oxide polymeric embryos and substrate surfaceswhich affect local saturations near the substrate surfacescontrolled heterogeneous growth rates. Water contact angles (7.5° ± 0.7, 22.8° ± 1.7, and 44.8° ± 3.7 for quartz, muscovite, and corundum, respectively) indicate that corundum has the highest substrate−water interfacial energy. Furthermore, a comparison of structural mismatches between the substrates and precipitates indicates a lowest precipitate− substrate interfacial energy for corundum. The fastest nucleation on corundum suggests that interfacial energies in the solution− substrate−precipitate system controlled heterogeneous nucleation rates. The unique information provided here bolsters our understanding of nanoparticle−mineral surface interactions, mineral surface modification by iron oxide coating, and pollutant transport.
1. INTRODUCTION
where k is the rate constant and f(Ω) is a function of the saturation ratio (Ω), which is the ratio between the actual dissolved compositions (Q) and the corresponding equilibrium compositions (K). The mineral surface area (A), which is a function of the number and size distribution of the precipitates, is initially created by nucleation and modified by both nucleation and growth. Because of limited experimental data on nucleation and growth kinetics, the mineral surface area is in general the least accurate parameter in reactive transport modeling.22 This inaccuracy significantly hinders better prediction of Fe(III) (hydr)oxide precipitation. In addition, the adsorption capacities of Fe(III) (hydr)oxide nanoparticles for pollutants are strongly size-dependent.23,24 Heterogeneous nucleation and growth rates of Fe(III) (hydr)oxide can affect both the sizes of the precipitates and their distributions among different mineral surfaces. Thus, to better predict pollutant transport, it is crucial to measure these rates on environmentally abundant mineral surfaces. Quartz (SiO2 ), muscovite (white mica, KAl2(Si 3Al)O10(OH,F)2), and corundum (Al2O3) are abundant minerals in the earth’s crust. The chosen crystal planes differ in surface
The heterogeneous nucleation and growth (i.e., growth of heterogeneously precipitated particles on mineral surfaces) of Fe(III) (hydr)oxides are important processes in natural and engineered aquatic systems. Often Fe(III) (hydr)oxides precipitate on other rocks’ surfaces via heterogeneous nucleation and change the reactivity of the pre-existing rocks.1−3 For example, it has been reported that iron oxide coated quartz has a much higher reactive surface area and adsorption capability for pollutants (e.g., As and heavy metals) than pure quartz.2,4,5 During heterogeneous Fe(III) (hydr)oxide precipitation, organic ligands,6 heavy metals (Cu, Zn, Pb, Cr, Ni, Cd, Mo, V, etc.),7−9 and arsenic10−15 can be immobilized on mineral surfaces through adsorption and coprecipitation processes.15−17 This immobilization process has been reported at acid mine drainage18−20 and managed aquifer recharge field sites,21 where arsenic mobilization from mineral dissolution (e.g., arsenopyrite, FeAsS) is a major challenge. To predict Fe(III) (hydr)oxide precipitation and the transport of aqueous pollutants in field sites, reactive transport models can be used. In these models, the rate of mineral precipitation from aqueous solution can be written as follows:22 Rate = Akf (Ω) = Akf (Q /K ) © 2013 American Chemical Society
Received: Revised: Accepted: Published:
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Figure 1. GISAXS scattering intensities cutting along the Yoneda wing from particles on quartz (A), mica (B), and corundum (C) surfaces immersed in 10−4 M Fe3+ solutions with 10 mM NaNO3. The black lines show the fitted curves. The in situ GISAXS/SAXS experiments were conducted at 20 °C for 1 h. Throughout the experiments, water evaporation was very slow, and no significant volume change of the original 1 mL solution occurred.
minerals.46 In addition, the ease of perfect cleavage along the muscovite (001) plane and its availability in high-grade natural and synthetic forms makes it a favorite substrate for epitaxially grown metal and crystal layers in diverse fields.29,30,35−42 Corundum is also an abundant mineral in the environment. In addition, the c-plane of corundum has been used for epitaxial growth of some oxides (e.g., ZnO) and metal (Fe) films.34 All the samples used had atomically flat surfaces (AFM images in Figure S2). All samples were cut into 1 cm × 1 cm squares to fit into the GISAXS/SAXS cell and were cleaned before measurements, as described in the Supporting Information. To prepare a 10−4 M Fe3+ solution with 10 mM ionic strength, we used reagent grade Fe(NO3)3, NaNO3, and ultrapure water. The measured pH values of the solution throughout the experiments were 3.7 ± 0.2. On the basis of Geochemist’s Workbench (GWB, Release 8.0, RockWare, Inc.) calculations, the initial solution was supersaturated with regard to Fe(OH)3 at 20 °C, with a saturation index (SI = log(Q/K) = log Ω) of 0.31. As Fe(OH)3 is a simple description of poorly crystallized hydrous iron (hydr)oxide rather than the real composition of the precipitates, there is a possibility that the actual saturation ratio of the precipitates is somewhat different from the value calculated here. However, the same aqueous solution was used equally for the three systems with different substrates in our study. Thus, our discussion and conclusions of the relative nucleation and growth rates of Fe(III) (hydr)oxides on different substrates are unaffected. 2.2. In Situ GISAXS Measurements. At 12 keV X-ray energy, an incident angle of 0.13° was chosen. At this incident angle, the X-ray probed only the structures at the substrate surfaces, and the calculated reflectivities of quartz, muscovite, and corundum were similar (with densities of 2.65, 2.82, and 3.99 g/cm3, the calculated reflectivities were 97.9%, 98.0%, and 98.7%, respectively).47,48 For each in situ run, a fresh 1 cm × 1 cm square substrate was placed inside a specially designed SAXS/GISAXS fluid cell,49 and the top surface of the substrate was aligned with the middle of the X-ray beam. The in situ run started as soon as 1 mL of freshly mixed solution was injected into the GISAXS cell. Only 2 min elapsed before the first GISAXS image was taken. During the run, time-resolved scattering signals were collected from particles precipitated on the substrate 1 frame/30 s by a 2D Pilatus 2 M detector
structures (details in Figure S1 in Supporting Information), pHiep values,25 and hydrophobic/hydrophilic properties. These surface properties can affect the interactions (e.g., electrostatic forces, adsorption, and bond formation) between these surfaces and the local aqueous species.26−28 Therefore, the local saturation ratio (Ω) near these substrate surfaces can differ from that of the bulk solution.26,29−31 The diverse surface properties of these minerals can also affect the interfacial energies within the solution−substrate−precipitate system.31,32 Thus, it is important to quantify the different heterogeneous nucleation and growth rates of Fe(III) (hydr)oxide on these surfaces and to identify the controlling mechanisms.22,33 Until now, the heterogeneous nucleation and growth rates of Fe(III) (hydr)oxide on these mineral surfaces have not been clearly quantified and compared due to the technical difficulty in probing the heterogeneously precipitated nanoparticles in situ at the mineral−water interface. In this study, we aimed to fill this important information gap. Our previously developed grazing incidence small-angle X-ray scattering (GISAXS) setup was employed to quantify in situ size and volume evolutions of Fe(III) (hydr)oxide nanoparticles precipitated at substrate−liquid interfaces26,31 and compare the heterogeneous nucleation and growth rates on different substrates. To elucidate the controlling mechanisms, the physicochemical and structural properties of the nanoparticles and the substrates were characterized with complementary techniques, including electrophoretic mobility, contact angle measurements, high resolution transmission electron microscopy (HRTEM), and X-ray diffraction (HRXRD).
2. MATERIALS AND METHODS 2.1. Substrate and Solution Preparation. We purchased synthetic single crystal quartz (Princeton Scientific Co.), muscovite (Goodfellow Cambridge Limited), and corundum (MTI Corporation). Quartz (102) surfaces, muscovite (001) surfaces, and corundum (0001) surfaces (C-plane) were chosen in this study based on their natural abundance and broad industrial applications.29,30,34−43 For example, quartz, very abundant in earth’s crust, has been widely used as the substrate of metal film and hematite nanoparticle formation.44,45 The muscovite (001) surface, abundant in natural environments, is structurally similar to the dominant surfaces of many clay 9199
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Figure 2. Evolutions of average radii (A) and the total volumes (B) and numbers (C) of the nanoparticles precipitated on the substrate surfaces from 10−4 M Fe3+ solutions with 10 mM NaNO3. In image A, the black lines show the linear regressions of the average particle sizes (Rg) over reaction time (t), the linear relationships are Rg_quartz = 0.6t + 2.6 (R2 = 0.96), Rg_mica = 0.6t + 1.3 (R2 = 0.91), and Rg_corundum = 0.4t + 1.4 (R2 = 0.93).
are shown in Figure 2, and the differences between the triplicate samples were within the error bars shown in Figure 2. 2.4. Atomic Force Microscopy, Zeta Potential (ζ), and Contact Angle Measurements. As a supplement to the in situ GISAXS measurements, ex situ atomic force microscopy (AFM) measurements were conducted after the GISAXS experiments to observe the precipitates on the substrate surfaces. AFM tapping mode (Veeco Inc.) was used, to collect height, amplitude, and phase images simultaneously. Probes were 125-μm-long with phosphorus (n) doped silicon tips (MPP-11100-10, Veeco probes). Images were collected with drive frequencies between 312 and 320 kHz, typical spring constants of 20−80 N/m, and a scan rate of 0.80 Hz. Nanoscope 7.20 software was used to analyze topographic features. A Zetasizer instrument (Nano ZS, Malvern Instruments Ltd.) was used for zeta potential and hydrodynamic particle size measurement. Freshly mixed solution was injected into a zeta cell (DTS1060C, Malvern Instruments), and the zeta potentials of the Fe(III) (hydr)oxide precipitates were measured every minute for 1 h at 20 °C. Because of the technical difficulty in measuring zeta potentials of single crystal surfaces, zeta potentials of quartz, mica, and corundum powders were measured under our experimental conditions (pH = 3.7 ± 0.2 in 10 mM NaNO3 solution). Water contact angles on the surfaces were measured in the air using a Phoenix 300 contact angle analyzer (Surface Electro Optics Co. Ltd., Korea). Each surface was placed on a horizontal holder and brought into contact with a water droplet hanging from a vertical needle connected to a syringe. Video acquisition was performed using a CCD camera, and 10 sequential measurements of the contact angles were conducted using a goniometer within 10 s of when the droplet touched the surface. Triplicate samples were measured. 2.5. Precipitate Phase Identification. To identify the phases of the precipitates on the substrates, first, Raman spectroscopy data were recorded using a Raman microscope (Renishaw, U.K.) with a 633 nm excitation wavelength; however, because the nanoparticle films on the substrates were too thin to generate enough signal or were poorly crystallized or amorphous, no peaks except those from the substrates were detected. Second, we took off the precipitates from the substrates by sonicating them in ethanol. Then, a drop of this suspension was placed on a Formvar/carbon-coated Cu grid and dried overnight in a desiccator. Electron diffraction
(Dectris Ltd., Baden, Switzerland). Meanwhile, incident and transmitted X-rays were detected by the ionic chamber (IC) and photodiode, respectively. The geometry of the GISAXS measurements and more details can be found in our previous studies.26,49 Experiments were conducted at beamline 12 ID-B at the Advanced Photon Source (APS), Argonne National Lab (ANL), IL. 2.3. GISAXS Data Analysis. First, the IC and photodiode values measured were examined carefully to confirm that the Xray beam intensity was stable and the alignments were good throughout all measurements. Thus, the differences in the scattering intensities came from the sample, not the X-ray beam’s condition. Second, for each experiment set, the first 2D scattering image was used as the background and was subtracted from each following 2D image; then, the subtracted 2D image was reduced to 1D by doing line-cuts along the Yoneda wing, where the scattering signal is enhanced along the in-plane direction due to the grazing incidence effect in GISAXS.50,51 Figure 1 shows the GISAXS scattering intensities (I) plotted against the scattering vector q (unit: Å −1) from particles precipitated on the substrate surfaces. Next, the GISAXS 1D scattering curves (Figure 1) were fit to the polydisperse sphere model with a structure factor included. Schultz distributions were used to get the particle size distributions, and the radii of gyration (Rg) of the particles were calculated (Figure 2A).52 The invariant Q,53 which is 2 defined as Q = ∫ ∞ 0 I(q)q dq and is proportional to the total volumes of precipitated nanoparticles (V, Figure 2B), was also calculated based on the fitted curves. The fitted function to the experimental data of intensity (I) over wave factor (q) allows extrapolating the intensity from 0 to infinity, and a converged value for the invariant can be obtained. Assuming a spherical particle shape and using the total particle volumes (V, Figure 2B) and the radii of gyration (Rg, Figure 2A), the total particle numbers (N, Figure 2C) were calculated in relative units: N = V/Rg3. GISAXS data reduction was performed with the GISAXS-SHOP macro, available at APS beamline 12 ID. All data analysis was performed with the Igor Pro program (V. 6.22A, Wave Metrics, Inc., Oregon). Triplicate experiments were run for quartz, mica, and corundum, and consistent scattering curves, in terms of both curve shape and scattering intensity, were obtained. Thus, the calculated particle sizes (in terms of radius of gyration, Rg) and total particle volumes were also consistent for triplicate experiments. The average values 9200
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measurements were performed for the nanoparticles on the grid using HRTEM (JEOL JEM-2100F field emission); however, no diffraction patterns were obtained, indicating that the precipitates were amorphous. To facilitate phase identification, the homogeneously precipitated particles in solution were collected by centrifuging with Millipore Amicon ultra-15 centrifugal filter units. The nanoparticles on filters were then transferred to glass slides and dried for 2 h in a desiccator. After a seven-day aging, HRXRD analysis was conducted at sector 11-BM of the APS at ANL, and broad peaks of ferrihydrite were observed.26 On the basis of all this information and our previous experiments with ex situ grazing incidence wide-angle X-ray scattering (where no diffraction peak was found),26,31 we concluded that the precipitates formed on the substrates were poorly crystallized Fe(III) (hydr)oxides, e.g., ferrihydrite.
Figure 3. Ex situ AFM observations of quartz (A), mica (B), and corundum (C) surfaces after reactions. The typical lateral resolution of the tapping mode used here was ∼40 nm, which is much bigger than the real particle sizes, while the vertical resolution was sub-Ångstrom. Thus, the real particle sizes were measured accurately in the vertical direction, which are quantified by the line cut curves below the images.
3. RESULTS 3.1. Size and Relative Total Volume Quantification of the Precipitates on Substrates. GISAXS scattering intensities from precipitates on different substrate surfaces are plotted (Figures 1A−C) against the scattering vector q (Å−1) at different reaction times. The shapes of the scattering intensity curves are related to the particle size and particle shape. The scattering intensities correlate to the total amounts of particles on the substrate surfaces. Interestingly, these curves show that for the same aqueous conditions, the heterogeneously precipitated Fe(III) (hydr)oxide nanoparticles on different mineral surfaces varied in terms of both size and total amount of particles, indicating that the heterogeneous nucleation and growth of Fe(III) (hydr)oxides were significantly influenced by substrate identity. As indicated by the orange arrows in Figure 1, as the reactions continued, the peak positions of the scattering intensity curves shifted to lower q. Because q is reciprocally related to particle size, these peak shifts indicate increases in particle sizes. To get more accurate particle sizes, the scattering curves (Figures 1) were fitted with the polydisperse sphere model with a structural factor included, and the average radii of gyration (Rg) of the particles were calculated (Figure 2A). Particle size fitting was not performed for quartz, mica, and corundum until 7, 20, and 9 min, respectively, because the low X-ray scattering intensities early in the reactions were insufficient to obtain good fitting. At the end of 1 h experiments, the in situ average radii (Rg) of the particles were 6.2 ± 0.3, 4.9 ± 0.5, and 3.8 ± 0.2 nm on quartz, mica, and corundum, respectively. The fitted average particle sizes (Figure 2A, Rg) on mica had a relatively large error, because the small amount of particles formed on mica generated low intensity scattering data. Ex situ AFM measurements conducted on dry samples showed 2−4 nm particles (Figures 3A−C), which were slightly smaller than the sizes measured by in situ GISAXS at the end of the 1 h experiments (4−6 nm). This difference could be due to dehydration of the ex situ samples. On the basis of invariant calculations from the fitted curves (Figure 1), the evolution of total particle volumes (V, in relative units) on the substrate surfaces are plotted over reaction time (t) in Figure 2B. The slopes of the curve, which represent increases in total particle volume per unit time and reflect both nucleation and growth, were defined as the precipitation rates. During the first 20 min, faster precipitation occurred on quartz than on corundum. Later on, faster precipitation occurred on corundum than on quartz. The slowest precipitation occurred on mica, where the data intensities were too low to analyze
until 20 min had elapsed. At the end of the 1 h experiments, the total particle volumes on quartz and corundum were similar and were around 4 times higher than on mica. 3.2. Surface Charges of the Precipitates and the Substrates. To understand the electrostatic interactions between the surfaces and the aqueous species, zeta potential measurements (ζ, Table 1) were conducted. The hydrous Table 1. Measured Zeta Potential (ζ) Values with Powder Samples and Reported pHiep of Different Substrates ζa (mV)
minerals Fe(III) (hydr) oxides
35.1 ± 4.0
quartz
−14.3 ± 2.6
muscovite corundum
−17.3 ± 4.2 41.7 ± 6.3
pHiepb 6.8−8.1(ferrihydrite), 8.5−9.5 (hematite) 2.0−3.0 always negative 8−10 (powder), 5−7 (crystal plane)
ref. 3, 9, 25, 54−56 3, 25, 57, 58 64, 87 25, 59−63
a
The zeta potential values of Fe(III) (hydr)oxides were measured every minute for 1 h. The signal was weak and not stable during approximately the first 20 min because of the small number of nanoparticles formed in solution. Later, the values were stable and are given here. The zeta potential values of quartz, mica, and corundum were measured with their powders under our experimental conditions (i.e., in 10 mM NaNO3 with pH adjusted to 3.7 ± 0.1 using HNO3). b The pHiep values are cited from references. The basal surface of muscovite is always negatively charged; thus, no pHiep value is applied.
Fe(III) oxide precipitates (ζ = 35.1 ± 4.0 mV) and the corundum powder (ζ = 41.7 ± 6.3 mV) were positively charged, while the quartz (ζ = −14.3 ± 2.6 mV) and mica (ζ = −17.3 ± 4.2 mV) powders were negatively charged. Our zeta potential measurements agree well with surface charges expected from the reported pHiep values of these minerals. As listed in Table 1, the pHiep values for iron oxides3,9,25,54−56 and quartz3,25,57,58 are reported as 6.8−9.5 and 2.0−3.0. The pHiep values for corundum are reported to be in the range of 8−10 for powders25,59,60 and 5−7 for the single crystal planes61−63 (Table 1). Thus, based on these reported pHiep values, at our experimental pH (3.7 ± 0.2), the iron oxides and corundum, which have pHiep values higher than the experimental pH condition, should be positively charged. 9201
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there, further growth can occur by the addition of embryos and monomers to the nuclei.33
Quartz, which has a pHiep value lower than the experimental pH conditions, should be negatively charged. For the (001) plane of muscovite, the Si/Al tetrahedral layer was exposed. Because of the isomorphic substitution of tetravalent silicon (Si4+) by trivalent aluminum (Al3+), this surface has a permanent negative structural charge,28,46,58 and its surface charge is independent of pH.28,64−66 Considering the potential differences in zeta potential values between powder samples and single crystal surfaces, we consulted the reported surface charges of these single crystal surfaces. For the (0001) plane of corundum, a zeta potential of 18 ± 3 mV at pH = 3.7 was measured using the streaming potential technique.67 Quartz single crystal surfaces were found to be negatively charged at pH 3.7, using both streaming potential measurement and XPS spectra calculation.68 These reported surface charges of quartz and corundum single crystals agree with our powder measurements in terms of being positive or negative; thus, we conclude that, under our experimental conditions, the quartz and mica surfaces were negatively charged, while the corundum surface was positively charged. However, there is a caveat that the absolute values of zeta potentials of these substrates measured with powders in this study can be different from those of single crystal surfaces. 3.3. Water Contact Angles on Substrates and Their Relation to Substrate−Water Interfacial Energies. To elucidate the role of the interfacial energy barrier on heterogeneous nucleation, we considered the substrate−water interfacial energies. However, there are only a few reported values in the literature, and these values are not very consistent with each other. For example, quartz−water interfacial energy (σquartz‑water) values of 360, 340, 120, and 416 mJ/m2 have been reported.69−72 Such inconsistency may be caused by the different treatment of these substrate surfaces or by different measuring methods, as the substrate−water interfacial energy can be affected by many factors, such as the surface roughness and coating, as well as the cleaning procedures. To obtain consistent data for comparison between substrate surfaces used in this study, water contact angles on these substrates were measured in the air. Substrates utilized were the same quality and were cleaned following the same procedures (refer to the Supporting Information) as in GISAXS experiments. The water contact angles of the substrates measured in the air were 7.5° ± 0.7, 22.8° ± 1.7, and 44.8° ± 3.7 for quartz, muscovite, and corundum, respectively. Less hydrophilic surfaces have larger water contact angles and would therefore also have higher substrate−water interfacial energies (σsubstrate‑water).73 Thus, we determined a semiquantitative ranking of these interfacial energies to be σquartz‑water < σmica-water < σcorundum‑water.
Hydrolysis: Fe3 + + 3H 2O → Fe(OH)3 + 3H+
(2)
Olation: 2[Fe(H 2O)5 OH2 +] → [(H 2O)4 Fe ‐ (OH)2 ‐ Fe(H 2O)4 ]4 + + 2H 2O
(3)
Oxolation: 2[Fe(H 2O)5 OH2 +] → [(H 2O)4 Fe ‐ O ‐ Fe(H 2O)5 ]4 + + 2H 2O
(4)
Once stable nuclei form, Fe(III) (hydr)oxide precipitation begins. From then on, as continuous nucleation and growth occur, the number and size of the precipitates will vary accordingly. Nucleation, which creates new stable nuclei from solution, increases the particle number while maintaining the particle size. Because particles are generated through nucleation only, the particle numbers thus represent the nucleation rates. On the other hand, during the growth process, Fe(III) (hydr)oxide polymeric embryos and monomers in solution attach to the existing particles. Therefore, the particle size increases while the particle number remains constant. As observed by AFM (Figure 3), the heterogeneously precipitated particles on the substrates are individual particles, rather than continuously grown films. Thus, the heterogeneous growth rate is defined as the increase in particle size per unit time. On the basis of these definitions, in the following sections 4.2 and 4.3, the heterogeneous nucleation and growth rates on the substrates are compared, and potential mechanisms are discussed. 4.2. Comparatively Slow Heterogeneous Growth on Corundum: Role of Electrostatic Forces. To compare the heterogeneous growth rates on the substrates, linear regressions (black lines in Figure 2A) were conducted for the average particle size (Rg) evolutions over reaction time (t). On the basis of the slopes of these regressions, the heterogeneous growth rates (unit: nm/min) of Fe(III) (hydr)oxide nanoparticles on mica and quartz were similar considering the error range, and both were faster than on corundum. The heterogeneous growth process requires the diffusion and attachment of polymeric embryos and monomers in solution to the existing particles on the substrate surfaces. Thus, the local saturation ratio (Ω = Q/K), which can be defined as the local concentrations of embryos and monomers (Q) divided by their equilibrium concentrations (K), is the driving force for heterogeneous growth. Higher concentrations of Fe(III) (hydr)oxide embryos and monomers in the local solution near the surfaces promote faster heterogeneous growth. Because the bulk solutions were identical for all experiments using different substrates, to explain the different heterogeneous growth rates, we evaluated the different substrate surface properties which could affect their local solution compositions. First, electrostatic forces between the substrate surfaces and the aqueous ions were considered. As shown in eq 2, the hydrolysis reactions produce Fe(III) (hydr)oxide monomers as the prerequisite for polymeric embryos formation. The hydrolysis reaction is promoted by high Fe3+ and low H+ concentrations. Therefore, to have high local monomer concentrations, high local Fe3+ and low H+ concentrations are preferred. As discussed in section 3.2, quartz and mica surfaces were similarly negatively charged, while corundum was positively charged. Considering this effect on the local
4. DISCUSSION 4.1. Defining the Nucleation and Growth Rates during Fe(III) (Hydr)oxide Precipitation. Before comparing the nucleation and growth rates on the substrates and exploring the mechanisms, we considered the pathways of Fe(III) (hydr)oxide precipitation. The first step is hydrolysis of ferric ion to form the Fe(OH)3 monomer (eq 2), followed by polymerization through continuous olation (hydroxo−bridging) and oxolation (oxo−bridging) reactions (eqs 3 and 4).74,75 When the size of the polymer is not big enough to be thermodynamically stable, it is called a polymeric embryo. Once the size of the polymer is larger than the critical size, it becomes thermodynamically stable and is termed a nucleus. From 9202
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distributions of aqueous ions, the aqueous cations (Na+, H+, and Fe3+) would be attracted toward the mica and quartz surfaces, while the anions (OH−, NO3−) would be attracted to the corundum surfaces. Considering only the electrostatic forces and not the surface adsorption of specific ions, a high local Fe3+ concentration requires a negatively charged surface, while a low H+ concentration requires a positively charged surface. Therefore, a high local Fe3+ and a low H+ concentration cannot be achieved at the same time for either positively or negatively charged surfaces. Thus, the effects of the electrostatic forces between the charged surfaces and aqueous cations and anions on the local Fe(III) (hydr)oxide monomer concentration offset each other. Second, we considered the electrostatic forces between the substrate surfaces and the Fe(III) (hydr)oxide polymeric embryos and monomers in solution, which can affect the local embryo and monomer concentrations near the substrate surfaces, thus affecting the local saturations. The Fe(OH)3 monomers are neutrally charged, so there will not be electrostatic forces between the monomers and the substrate surfaces. For the Fe(III) (hydr)oxide embryos in solution, if we assume they were as positively charged as the Fe(III) (hydr)oxide particles (Table 1), there should be similar electrostatic attractive forces between the embryos in solution and the quartz and mica surfaces, while electrostatic repulsive forces exist between the embryos and the corundum surface. Given this, the slowest growth should occur on corundum, and similar growth rates should occur on quartz and mica surfaces, as observed in our experiments. This agreement suggests that the electrostatic forces between the charged embryos and the surfaces, which affect the local saturations, control the heterogeneous growth rates. 4.3. Comparatively Fast Heterogeneous Nucleation on Corundum: Role of Interfacial Energies. The calculated particle numbers (Figure 2C) on the substrates were used to compare the relative heterogeneous nucleation rates. The particle numbers on corundum increased with time (indicated by the guideline in Figure 2C) and were the most among the three substrates, indicating the fastest nucleation on corundum. For quartz and mica, as there is some overlap in their particle numbers throughout the reactions, no comparisons were made regarding their nucleation rates. To explore the possible mechanisms causing the fastest heterogeneous nucleation to occur on corundum despite its positive charge, we recall classical nucleation theory (CNT). According to classical nucleation theory,33 for heterogeneous nucleation to occur, a free energy barrier (ΔG*) must be overcome. Taking the nucleation barrier as the activation energy for nucleation to occur, the nucleation rate (I) can be described in eq 5, where k is the rate constant: ⎛ ΔG* ⎞ −3 −1 ⎟ cm I = k exp⎜ − s ⎝ RT ⎠
where Apl and Aps are the surface areas between the precipitate−liquid and precipitate−substrate interfaces.33 Considering both the interfacial energy (ΔGinterface) and the energy barrier for bulk crystal growth (ΔGr), a higher saturation ratio (Ω = Q/K) and substrate−solution interfacial energy (σsl) and a lower precipitate−substrate interfacial energy (σps) are preferred for faster heterogeneous nucleation.33 The first parameter, the local saturation (Ω), cannot be the dominant mechanism controlling the heterogeneous nucleation, as the local saturation index (Ω) was expected to be the lowest on corundum, as discussed in section 4.2. We then considered the substrate−solution interfacial energy (σsl). As discussed in section 3.3, based on water contact angle measurements, the order of the σsl values was σquartz−water < σmica−water < σcorundum−water. On the basis of eqs 5 and 6, a higher substrate−solution interfacial energy (σsl) can result in a lower energy barrier for heterogeneous nucleation and hence promotes heterogeneous nucleation; thus, similarly, the highest corundum−water interfacial energy results in the fastest heterogeneous nucleation rate on corundum. Finally, we considered the precipitate−substrate interfacial energy (σps), which is related to the crystal structural mismatch between the substrate and the precipitates.76−80 A higher structural mismatch results in a higher σps.76−80 Because the precipitates and the substrates belong to different crystal systems (Table S1 in Supporting Information),47 we cannot calculate the exact lattice mismatch, but the metal−oxygen and oxygen−oxygen bond length mismatches can be compared as an alternative.77 As shown in Table S2,47,81,82 among the three substrates, the bond length mismatch is the smallest between corundum and the precipitates; i.e., the Fe−O (2.0 Å) and O− O (2.8 Å) bond lengths of Fe(III) (hydr)oxides are similar to the Al−O (1.9 Å) and O−O (2.7 Å) bond lengths of corundum.47,81,82 The smallest crystal structural mismatch between corundum and the precipitates should therefore result in the lowest precipitate−substrate interfacial energy for corundum, promoting the fastest heterogeneous nucleation on corundum compared to the other two substrates. To sum up, our observations and analysis suggested that the interfacial energies (σsl and σps) within the liquid−substrate− precipitate system played a more crucial role in the heterogeneous nucleation process than the local saturations, which controlled the heterogeneous growth rates. Considering the importance of heterogeneous nucleation and the extremely limited data currently available for the interfacial energies of the liquid−substrate−precipitate system, further exploration of this area is an important future direction. To more accurately calculate the interfacial energies between the substrate surfaces and the precipitates, we could measure the in situ contact angles between the substrate surfaces and the precipitates,49 or study the heterogeneous nucleation rates under different saturation indices.32 Further exploration in this area is our undergoing research and is beyond the scope of the current manuscript.
(5)
5. ENVIRONMENTAL IMPLICATIONS Using GISAXS, this study provides for the first time quantitative information about both the size and volume evolutions of Fe(III) (hydr)oxide particles precipitated on environmentally ubiquitous mineral surfaces. This technique allowed us to separately look into the heterogeneous nucleation and growth processes. Interestingly, both the fastest nucleation and slowest growth occurred on corundum. Interfacial energies and local saturations were found to be the controlling
The energy barrier of heterogeneous nucleation originates from both the energy barrier for bulk crystal growth (ΔGr = −RT ln Ω) and the interfacial free energy (ΔGinterface), which originates from the creation of the precipitate (p)−substrate (s) interface and the precipitate (p)−solution (l) interface, as well as the partial covering of the substrate (s)−solution (l) interface, as shown in eq 6: ΔG interface=A pl σpl + A ps(σps − σsl)
(6) 9203
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mechanisms for heterogeneous nucleation and growth, respectively. In addition, this study provides valuable insights into the heterogeneous nucleation and growth of Fe(III) (hydr)oxides that can be incorporated into reactive transport modeling. Although current reactive transport models have significantly improved our ability to predict the fate and transport of contaminants, there are still limitations in providing accurate data on the rates of heterogeneous nucleation and growth and resulting changes in mineral surface area and reactivity. To address these limitations, we provided new quantitative information that can help predict contaminant transport and nutrient availability in soils and groundwater environments. Heterogeneous nucleation and growth from solution are also important processes in materials science, e.g., particle size control in thin film development.83−85 Here, we used GISAXS to separate the heterogeneous nucleation and growth rates, and together with supplementary analysis of the surface and nanoparticle properties, to elucidate possible mechanisms. Quartz, muscovite, and corundum have been widely used in industrial applications.29,30,34−42,86 The fundamental knowledge of different surface processes gained in this study can be valuable for environmentally benign materials synthesis (e.g., heterogeneous catalyst synthesis, nanoparticle deposition, and semiconductor processing).
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REFERENCES
(1) Scheidegger, A.; Borkovec, M.; Sticher, H. Coating of silica sand with goethite: preparation and analytical identification. Geoderma 1993, 58 (1−2), 43−65. (2) Rusch, B.; Hanna, K.; Humbert, B. Coating of quartz silica with iron oxides: Characterization and surface reactivity of iron coating phases. Colloids Surf., A 2010, 353 (2−3), 172−180. (3) Abudalo, R. A.; Bogatsu, Y. G.; Ryan, J. N.; Harvey, R. W.; Metge, D. W.; Elimelech, M. Effect of ferric oxyhydroxide grain coatings on the transport of bacteriophage PRD1 and Cryptosporidium parvumoocysts in saturated porous media. Environ. Sci. Technol. 2005, 39, 6412−6419. (4) Wu, K.; Liu, R.; Liu, H.; Zhao, X.; Qu, J. Arsenic (III, V) adsorption on iron-oxide-coated manganese sand and quartz sand: comparison of different carriers and adsorption capacities. Environ. Eng. Sci. 2011, 28 (9), 643−651. (5) Ying, X.; Axe, L. Synthesis and characterization of iron oxidecoated silica and its effect on metal adsorption. J. Colloid Interface Sci. 2005, 282 (1), 11−19. (6) Boukhalfa, C. Sulfate removal from aqueous solutions by hydrous iron oxide in the presence of heavy metals and competitive anions Macroscopic and spectroscopic analyses. Desalination 2010, 250 (1), 428−432. (7) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurences and Uses; Wiley-VCH: Darmstadt, Germany, 2003. (8) Potter, H. A. B.; Yong, R. N. Influence of iron/aluminium ratio on the retention of lead and copper by amorphous iron-aluminium oxides. Appl. Clay Sci. 1999, 14 (1−3), 1−26. (9) Brinza, L.; Benning, L. G.; Satham, P. J. Adsorption studies of Mo and V onto ferrihydrite. Mineral. Mag. 2008, 72 (1), 385−388. (10) Jia, Y.; Demopoulos, G. P. Coprecipitation of arsenate with iron(III) in aqueous sulfate media: Effect of time, lime as base and coions on arsenic retention. Water Res. 2008, 42 (3), 661−668. (11) Silva, J.; Mello, J. W. V.; Gasparon, M.; Abrahao, W. A. P.; Ciminelli, V. S. T.; Jong, T. The role of Al-Goethites on arsenate mobility. Water Res. 2010, 44 (19), 5684−5692. (12) Violante, A.; Gaudio, S. D.; Pigna, M.; Ricciardella, M.; Banerjee, D. Coprecipitation of Arsenate with Metal Oxides. 2. Nature, Mineralogy, and Reactivity of Iron(III) Precipitates. Environ. Sci. Technol. 2007, 41 (24), 8275−8280. (13) Violante, A.; Pigna, M.; Del Gaudio, S.; Cozzolino, V.; Banerjee, D. Coprecipitation of arsenate with metal oxides. 3. nature, mineralogy, and reactivity of iron(III) aluminum precipitates. Environ. Sci. Technol. 2009, 43 (5), 1515−1521. (14) Violante, A.; Ricciardella, M.; Del Gaudio, S.; Pigna, M. Coprecipitation of arsenate with metal oxides: nature, Mineralogy, and reactivity of aluminum precipitates. Environ. Sci. Technol. 2006, 40 (16), 4961−4967. (15) Fuller, C. C.; Davis, J. A.; Waychunas, G. A. Surface chemistry of ferrihydrite: Part 2. Kinetics of arsenate adsorption and coprecipitation. Geochim. Cosmochim. Acta 1993, 57 (10), 2271−2282. (16) Auffan, M.; Rose, J.; Bottero, J.-Y.; Lowry, G. V.; Jolivet, J.-P.; Wiesner, M. R. Towards a definition of inorganic nanoparticles from an environmental, health and safety perspective. Nat. Nanotechnol. 2009, 4, 634−641. (17) Grundl, T.; Delwiche, J. Kinetics of ferric oxyhydroxide precipitation. J. Contam. Hydrol. 1993, 14, 71−97. (18) Carlson, L.; Bigham, J. M.; Schwertmann, U.; Kyek, A.; Wagner, F. Scavenging of As from acid mine drainage by schwertmannite and ferrihydrite: A comparison with synthetic analogues. Environ. Sci. Technol. 2002, 36 (8), 1712−1719. (19) Tokoro, C.; Yatsugi, Y.; Koga, H.; Owada, S. Sorption mechanisms of arsenate during coprecipitation with ferrihydrite in aqueous solution. Environ. Sci. Technol. 2010, 44 (2), 638−643. (20) Zhu, M.; Legg, B.; Zhang, H.; Gilbert, B.; Ren, Y.; Banfield, J. F.; Waychunas, G. A. Early stage formation of iron oxyhydroxides during neutralization of simulated acid mine drainage Solutions. Environ. Sci. Technol. 2012, 46 (15), 8140−8147.
ASSOCIATED CONTENT
S Supporting Information *
Includes the experimental operation and X-ray scattering analysis details, substrate structural plots and AFM images (Tables S1 and S2, Figures S1 and S2), experimental setup (Figure S3), 2D GISAXS scattering images (Figure S4), and HRTEM (Figure S5) and HRXRD (Figure S6) results. This material is available free of charge via the Internet at http:// pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Phone: (314) 935-4539 Fax: (314) 935-7211. E-mail: ysjun@ seas.wustl.edu. Web: http://encl.engineering.wustl.edu/. Present Address §
Department of Civil and Environmental Engineering, University of Houston, Houston, TX 77204.
Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work is supported by a Washington University Faculty Startup Grant and the National Science Foundation’s Environmental Chemical Science Program (CHE-1214090). We thank Ms. Jessica Ray for HRTEM measurements and Dr. Soenke Seifert for beamline experimental help. We also thank Dr. Jill Pasteris for her insight in crystallography and Dr. Alejandro Fernandez-Martinez and Dr. Glenn W. Waychunas for valuable discussion about electron density and surface energy. Use of the Advanced Photon Source (Sector 11-BM and 12ID-B) at Argonne National Laboratory was supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, under Contract No. DE-AC02-06CH11357. 9204
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formation on phlogopite surfaces at 75−95°C and 75 atm. Energy Environ. Sci. 2012, 5, 5758−5767. (42) Hu, Y.; Jun, Y.-S. Biotite dissolution in brine at varied temperatures and CO2 pressures: Its activation energy and potential CO2 intercalation. Langmuir 2012, 28 (41), 14633−14641. (43) Eng, P. J.; Trainor, T. P.; Brown, G. E.; Waychunas, G. A.; Newville, M.; Sutton, S. R.; Rivers, M. L. Structure of the hydrated alpha-Al2O3 (0001) surface. Science 2000, 288 (5468), 1029−1033. (44) Uribe, J. D.; Osorio, J.; Barrero, C. A.; Girata, D.; Morales, A. L.; Devia, A.; Gomez, M. E.; Ramirez, J. G.; Gancedo, J. R. Hematite thin films: growth and characterization. Hyperfine Interact. 2006, 169 (1− 3), 1355−1362. (45) Glasscock, J. A.; Barnes, P. R. F.; Plumb, I. C.; Savvides, N. Enhancement of photoelectrochemical hydrogen production from hematite thin films by the introduction of Ti and Si. J. Phys. Chem. C 2007, 111 (44), 16477−16488. (46) Sposito, G.; Skipper, N. T.; Sutton, R.; Park, S. H.; Soper, A. K.; Greathouse, J. A. Surface geochemistry of the clay minerals. Proc. Natl. Acad. Sci. U. S. A. 1999, 96 (7), 3358−3364. (47) Chichagov, A. V. Information-Calculating System on Crystal Structure Data of Minerals (MINCRYST). Kristallographiya 1990, 35, 610−616. (48) Gullikson, E. X-ray interactions with matter: mirror reflectivity. http://henke.lbl.gov/optical_constants/mirror2.html. (49) Jun, Y. S.; Lee, B.; Waychunas, G. A. In situ observations of nanoparticle early development kinetics at mineral-water interfaces. Environ. Sci. Technol. 2010, 44 (21), 8182−8189. (50) Lee, B.; Seifert, S.; Riley, S. J.; Tikhonov, G.; Tomczyk, N. A.; Vajda, S.; Winans, R. E. Anomalous grazing incidence small-angle x-ray scattering studies of platinum nanoparticles formed by cluster deposition. J. Chem. Phys. 2005, 123 (7), 074701. (51) Renaud, G.; Lazzari, R.; Revenant, C.; Barbier, A.; Noblet, M.; Ulrich, O.; Leroy, F.; Jupille, J.; Borensztein, Y.; Henry, C. R.; Deville, J. P.; Scheurer, F.; Mane-Mane, J.; Fruchart, O. Real-time monitoring of growing nanoparticles. Science 2003, 300 (5624), 1416−1419. (52) Aragon, S. R.; Pecora, R. Theory of dynamic light scattering from polydisperse systems. J. Chem. Phys. 1976, 64 (6), 2395−2404. (53) Li, Y. J.; Gao, T.; Chu, B. Synchrotron SAXS studies of the phase-separation kinetics in a segmented polyurethane. Macromolecules 1992, 25 (6), 1737−1742. (54) Appelo, C. A. J.; Van der Weiden, M. J. J.; Tournassat, C.; Charlet, L. Surface complexation of ferrous iron and carbonate on ferrihydrite and the mobilization of arsenic. Environ. Sci. Technol. 2002, 36 (14), 3096−3103. (55) Echigo, T.; Aruguete, D. M.; Murayama, M.; Hochella, M. F., Jr. Influence of size, morphology, surface structure, and aggregation state on reductive dissolution of hematite nanoparticles with ascorbic acid. Geochim. Cosmochim. Acta 2012, 90, 149−162. (56) Gilbert, B.; Frandsen, C.; Maxey, E. R.; Sherman, D. M. Bandgap measurements of bulk and nanoscale hematite by soft x-ray spectroscopy. Phys. Rev. B 2009, 79, 3. (57) Churchill, H.; Teng, H.; Hazen, R. M. Correlation of pHdependent surface interaction forces to amino acid adsorption: implications for the origin of life. Am. Mineral. 2004, 89, 1048−1055. (58) Zaia, D. A. A review of adsorption of amino acids on minerals: was it important for the origin of life? Amino Acids 2004, 27, 113−118. (59) Sung, J.; Zhang, L.; Tian, C.; Shen, Y. R.; Waychunas, G. A. Effect of pH on the water/α-Al2O3 (1102) interface structure studied by sum-frequency vibrational spectroscopy. J. Phys. Chem. C 2011, 115 (28), 13887−13893. (60) Kershner, R. J.; Bullard, J. W.; Cima, M. J. Zeta potential orientation dependence of sapphire substrates. Langmuir 2004, 20 (10), 4101−4108. (61) Kosmulski, M. The pH-dependent surface charging and the points of zero charge. J. Colloid Interface Sci. 2002, 253 (1), 77−87. (62) Kosmulski, M. pH-dependent surface charging and points of zero charge III. Update. J. Colloid Interface Sci. 2006, 298 (2), 730− 741.
(21) Neil, C. W.; Yang, Y. J.; Jun, Y.-S. Arsenic mobilization and attenuation by mineral-water interactions: implications for managed aquifer recharge. J. Environ. Monit. 2012, 14 (7), 1772−1788. (22) Steefel, C.; Cappellen, P. v. A new kinetic approach to modeling water-rock interaction: The role of nucleation, precursors, and Ostwald ripening. Geochim. Cosmochim. Acta 1990, 54 (10), 2657− 2677. (23) Navrotsky, A.; Mazeina, L.; Majzlan, J. Size-driven structural and thermodynamic complexity in iron oxides. Science 2008, 319 (5870), 1635−1638. (24) Yean, S.; Cong, L.; Yavuz, C. T.; Mayo, J. T.; Yu, W. W.; Falkner, J. C.; Kan, A. T.; Colvin, V. L.; Tomson, M. B. Effect of magnetite particle size on adsorption and desorption of arsenite and arsenate. J. Mater. Res. 2006, 21 (7), 1862−1862. (25) Zhang, T. C.; Hu, Z.; Surampalli, R. Y.; Tyagi, R. D.; Lai, K. C. K.; Lo, I. M. C. Nanotechnologies for Water Environment Applications; American Society of Civil Engineers: Reston, VA, 2009. (26) Hu, Y.; Lee, B.; Bell, C.; Jun, Y.-S. Environmentally abundant anions influence the nucleation, growth, ostwald ripening, and aggregation of hydrous Fe(III) oxides. Langmuir 2012, 28 (20), 7737−7746. (27) Elhadj, S.; DeYoreo, J. J.; Hoyer, J. R.; Dove, P. M. Role of molecular charge and hydrophilicity in regulating the kinetics of crystal growth. Proc. Natl. Acad. Sci. U. S. A. 2006, 103 (51), 19237−19242. (28) Lee, S. S.; Fenter, P.; Nagy, K. L.; Sturchio, N. C. Monovalent ion adsorption at the muscovite (001)-solution interface: relationships among ion coverage and speciation, interfacial water structure, and substrate relaxation. Langmuir 2012, 28 (23), 8637−8650. (29) Hu, Y.; Ray, J. R.; Jun, Y. S. Biotite-brine Interactions under acidic hydrothermal conditions: fibrous Illite, goethite, and kaolinite formation and biotite surface cracking. Environ. Sci. Technol. 2011, 45 (14), 6175−6180. (30) Hu, Y.; Ray, J. R.; Jun, Y.-S. Na+, Ca2+, and Mg2+ in brines affect supercritical CO2-brine-biotite interactions: ion exchange, biotite dissolution, and Illite precipitation. Environ. Sci. Technol. 2013, 47 (1), 191−197. (31) Ray, J. R.; Lee, B.; Baltrusaitis, J.; Jun, Y.-S. Formation of Iron(III) (Hydr)oxides on Polyaspartate- and Alginate-Coated Substrates: Effects of Coating Hydrophilicity and Functional Group. Environ. Sci. Technol. 2012, 46 (24), 13167−13175. (32) Fernandez-Martinez, A.; Hu, Y.; Lee, B.; Jun, Y.-S.; Waychunas, G. A. In situ determination of interfacial energies between heterogeneously nucleated CaCO3 and quartz substrates: thermodynamics of CO2 mineral trapping. Environ. Sci. Technol. 2013, 47 (1), 102−109. (33) Lasaga, A. C. Kinetic Theory in the Earth Sciences; Princeton University Press: Princeton, NJ, 1998. (34) Dobrovinskaya, E. R.; Lytvynov, L. A.; Pishchik, V. Sapphire: Material, Manufacturing, Applications; Springer: New York, 2009. (35) Dishner, M. H.; Ivey, M. M.; Gorer, S.; Hemminger, J. C.; Feher, F. J. Preparation of gold thin films by epitaxial growth on mica and the effect of flame annealing. J. Vac. Sci. Technol., A 1998, 16 (6), 3295−3300. (36) Baski, A. A.; Fuchs, H. Epitaxial growth of silver on mica as studied by AFM and STM. Surf. Sci. 1994, 313 (3), 275−288. (37) Sun, L.-H.; Xu, C.-Y.; Yu, F.; Tao, S.-X.; Li, J.; Zhou, H.; Huang, S.; Tang, L.; Hu, J.; He, J.-H. Epitaxial Growth of Trichosanthin Protein Crystals on Mica Surface. Cryst. Growth Des. 2010, 10 (6), 2766−2769. (38) Utama, M. I. B.; Peng, Z.; Chen, R.; Peng, B.; Xu, X.; Dong, Y.; Wong, L. M.; Wang, S.; Sun, H.; Xiong, Q. Vertically aligned cadmium chalcogenide nanowire arrays on muscovitemica: A demonstration of epitaxial growth strategy. Nano Lett. 2010, 11 (8), 3051−3057. (39) Toman, K.; Taylor, P. R. Epitaxial growth of dolomite on mica. Am. Mineral. 1974, 59 (7−8), 871−872. (40) Grunbaum, E.; Schwarz, R. Epitaxial growth of titanium films on mica. J. Appl. Phys. 1969, 40 (8), 3364−9. (41) Garcia, D.; Shao, H.; Hu, Y.; Ray, J. R.; Jun, Y.-S. Supercritical CO2-brine induced dissolution, swelling, and secondary mineral 9205
dx.doi.org/10.1021/es401160g | Environ. Sci. Technol. 2013, 47, 9198−9206
Environmental Science & Technology
Article
(87) Bourg, I. C.; Sposito, G. Ion exchange phenomena. In Handbook of Soil Science, Properties and Processes; Huang, P. M., Li, Y., Sumner, M. E., Eds.; CRC Press: Boca Raton, FL, 2011; Vol. 16.
(63) Kosmulski, M. pH-dependent surface charging and points of zero charge - II. Update. J. Colloid Interface Sci. 2004, 275 (1), 214− 224. (64) Schlegel, M. L.; Nagy, K. L.; Fenter, P.; Cheng, L.; Sturchio, N. C.; Jacobsen, S. D. Cation sorption on the muscovite (001) surface in chloride solutions using high-resolution X-ray reflectivity. Geochim. Cosmochim. Acta 2006, 70 (14), 3549−3565. (65) Raviv, U.; Laurat, P.; Klein, J. Fluidity of water confined to subnanometre films. Nature 2001, 413 (6851), 51−54. (66) Raviv, U.; Klein, J. Fluidity of bound hydration layers. Science 2002, 297 (5586), 1540−1543. (67) Kershner, R. J.; Bullard, J. W.; Cima, M. J. Zeta potential orientation dependence of sapphire substrates. Langmuir 2004, 20, 4101−4108. (68) Duval, Y.; Mielczarski, J. A.; Pokrovsky, O. S.; Mielczarski, E.; Ehrhardt, J. J. Evidence of the existence of three types of species at the quartz-aqueous solution interface at pH 0−10: XPS surface group quantification and surface complexation modeling. J. Phys. Chem. B 2002, 106 (11), 2937−2945. (69) Parks, G. Surface and interfacial free energies of quartz. J. Geophys. Res. 1984, 89, 3997−4008. (70) Mizele, J.; Dandurand, J.; Schott, J. Determination of the surface energy of amorphous silica from solubility measurements in micropore. Surf. Sci. 1985, 162 (1−3), 830−837. (71) Rimstidt, J. D.; Cole, D. R. Geothermal mineralization. 1. the mechanism of formation of the beowawe, Nevada, siliceous sinter deposit. Am. J. Sci. 1983, 283 (8), 861−875. (72) Iler, R. K. The Chemistry of Silica: Solubility, Polymerization, Colloid and Surface Properties and Biochemistry of Silica; Wiley: New York, 1979; p 365. (73) Dong, J.; Mao, G.; Hill, R. M. Nanoscale aggregate structures of trisiloxane surfactants at the solid-liquid interface. Langmuir 2004, 20 (7), 2695−2700. (74) Flynn, C. M. Hydrolysis of inorganic iron(III) salts. Chem. Rev. 1984, 84 (1), 31−41. (75) Sharma, S. K. Adsorptive Iron Removal from Groundwater; Swets & Zeitlinger B.V.: Lisse, Holland, 2001. (76) Markov, I. V. Crystal Growth for Beginners: Fundamentals of Nucleation, Crystal Growth and Epitaxy; World Scientific Publishing Co. Pte. Ltd.: Singapore, 2003. (77) Jun, Y. S.; Kendall, T. A.; Martin, S. T.; Friend, C. M.; Vlassak, J. J. Heteroepitaxial nucleation and oriented growth of manganese oxide islands on carbonate minerals under aqueous conditions. Environ. Sci. Technol. 2005, 39 (5), 1239−1249. (78) Zheleva, T.; Jagannadham, K.; Narayan, J. Epitaxial growth in large lattice mismatch systems. J. Appl. Phys. 1994, 75 (2), 860−871. (79) Verlinden, B.; Driver, J.; Samajdar, I.; Doherty, R. D. ThermoMechanical Processing of Metallic Materials; Elsevier: Oxford, 2007. (80) Van der merwe, J. H.; Ball, C. A. B., Energy of interfaces between crystals. In Epitaxial Growth Part B; Matthews, J. W., Ed.; Academic Press: New York, 1975. (81) Xu, W.; Hausner, D. B.; Harrington, R.; Lee, P. L.; Strongin, D. R.; Parise, J. B. structural water in ferrihydrite and constraints this provides on possible structure models. Am. Mineral. 2011, 96, 513− 520. (82) Manceau, A. Critical evaluation of the revised akdalaite model for ferrihydrite. Am. Mineral. 2011, 96, 521−533. (83) Ramaswamy, V.; Haynes, T. E.; White, C. W.; MoberlyChan, W. J.; Roorda, S.; Aziz, M. J. Synthesis of nearly monodisperse embedded nanoparticles by separating nucleation and growth in ion implantation. Nano Lett. 2005, 5 (2), 373−377. (84) Chayen, N. E. Methods for separating nucleation and growth in protein crystallisation. Prog. Biophys. Mol. Biol. 2005, 88 (3), 329−337. (85) Saridakis, E.; Dierks, K.; Moreno, A.; Dieckmann, M. W. M.; Chayen, N. E. Separating nucleation and growth in protein crystallization using dynamic light scattering. Acta Crystallogr., Sect. D 2002, 58, 1597−1600. (86) Klein, C. The Manual of Mineral Science; John Wiley & Sons: New York, 2002. 9206
dx.doi.org/10.1021/es401160g | Environ. Sci. Technol. 2013, 47, 9198−9206
